Electrochemistry & Redox. Voltaic Cells. Electrochemical Cells

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Garry Dickerson
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Electrochemistry & Redox An oxidation-reduction (redox) reaction involves the transfer of electrons from the reducing agent to the oxidising agent.

Voltaic Cells A voltaic electrochemical cell involves two half cells one containing an oxidising agent and the other a reducing agent.

The PULL or DRIVING FORCE on the electrons is the cell potential (E cell ) or the electromotive force (emf) of the cell, measured in volts.

1 Electrochemistry & Redox An oxidation-reduction (redox) reaction involves the transfer of electrons from the reducing agent to the oxidising agent. OXIDATION - is the LOSS of electrons REDUCTION - is the GAIN of electrons Zn(s) Zn 2+ (aq) + 2e - (aq) Cu 2+ (aq) + 2e - (aq) Cu(s) These represents the redox HALF-EQUATIONS 1 Voltaic Cells A voltaic electrochemical cell involves two half cells one containing an oxidising agent and the other a reducing agent. These cells are connected with a wire, to allow electron flow and a salt bridge to complete the circuit and maintain electrical neutrality. The PULL or DRIVING FORCE on the electrons is the cell potential (E cell ) or the electromotive force (emf) of the cell, measured in volts. 2 Electrochemical Cells Voltaic (Galvanic) cells are those in which spontaneous chemical reactions produce electricity and supply it to other circuits. G < 0 An ox, red cat Anode: oxidation Reduction: cathode Electrolytic cells are those in which electrical energy causes non-spontaneous chemical reactions to occur. G > 0 3 1

2 Balancing Redox Equations The concept of Oxidation Number is artificial. In simple ions it is equivalent to the charge on the ion. Oxidation involves an increase in oxidation number Reduction involves a decrease in oxidation number As half cells, determine oxidation numbers and balance electrons. Combine half cells balancing gain/loss of electrons. Balance with H 2 O and H + or H 2 O and OH -. Check charges balance. 0 +II Zn(s) Zn 2+ (aq) + 2e - (aq) +II 0 Cu 2+ (aq) + 2e - (aq) Cu(s) Zn(s) + Cu 2+ (aq) Zn 2+ (aq) + Cu(s) 4 Standard Reduction Potentials In data tables half cells are written as reductions. Standard hydrogen electrode defined as E = 0 V (1 atm H 2, [H + ] = 1 M, at all temperatures). The more negative E, the greater the tendency to release electrons (act as a reducing agent). Fe 3+ + e - Fe 2+ E = 0.77 V Cu e - Cu E = 0.34 V 2H + + 2e - H 2 (g) E = 0.00 V Zn e - Zn E = V 5 Calculating Cell Potential When combining two half-reactions: One half-cell reaction is reversed (thus the sign of the reduction potential is reversed). Number of electrons lost must equal the number gained. Note cell potential is an INTENSIVE PROPERTY i.e. when a half- reaction is multiplied by and integer E stays the SAME. E.g. Fe 3+ + e - Fe 2+ E = 0.77 V 1 Cu e - Cu E = 0.34 V 2 reverse 2 Cu Cu e - E = V double 1 2Fe e - 2Fe 2+ E = 0.77 V add equations 2Fe 3+ + Cu 2Fe 2+ + Cu 2+ E cell = 0.77+(-0.34) = 0.43 V 6 2

3 Question Write a balanced equation for the oxidation of lactate to pyruvate and calculate the cell potential. O 2 + 4H + + 4e - 2H 2 O E = 0.82 V pyruvate +2H + +2e - lactate E =-019V G = -nfe max Example: Calculate G for a cell reaction: Cu 2+ (aq) + Fe(s) Cu(s) + Fe 2+ (aq) Is this a spontaneous reaction? Cu 2+ (aq) + 2e - Cu(s) Fe 2+ (aq) + 2e - Fe(s) E = 0.34 V E = V Cu 2+ (aq) + Fe(s) Cu(s) + Fe 2+ (aq) E cell = (+0.44) = 0.78 V G = - 2 x x 0.78 (F=96485 C mol -1 ) = x 10 5 J This process is spontaneous as indicated by the negative sign of G and the positive sign for E cell. 8 Glucose metabolism C 6 H 12 O 6 + 6O 2 6CO 2 + 6H 2 O G = kj mol -1 If all this energy was released at once it would totally swamp the cell. Instead energy in food used to create an electrochemical potential which is used to create the high energy molecule ATP. The energy can be released in small steps rather than all at once. 9 3

Metabolism and energy storage Redox reactions involved form part of the electron-transport chain (ETC). The ETC lies on the inner membrane of mitochondria.

4 Metabolism and energy storage Redox reactions involved form part of the electron-transport chain (ETC). The ETC lies on the inner membrane of mitochondria. The ETC involves a series of large molecules, mostly proteins, that use Fe 2+ /Fe 3+ to pass electrons down the chain. 10 Electron transport chain ETC The reaction that ultimately powers ETC is H 2 + ½O 2 H 2 O but do not have H 2 (g) but rather 2H + + 2e - in the presence of NADH. 2H + + 2e - + ½O 2 H 2 O NADH + H + NAD + + 2H + + 2e - NADH(aq) + H + (aq) + ½O 2 (aq) NAD + (aq) + H 2 O(l) E overall = 1.13 V and G = -218 kj mol -1. At each of the points shown there is enough free energy released to drive ADP 3- + HPO H + ATP 4- + H 2 O G = 30.5 kj mol Concentration Cells Imagine a cell: Cu(s) Cu 2+ (aq) Cu 2+ (aq) Cu(s) Low [Cu 2+ ] High [Cu 2+ ] Cu(s) Cu 2+ (aq) + 2e - Cu 2+ (aq) + 2e - Cu(s) 12 4

Concentration Cells Nerve cells operate as concentration cells. Inside cell [Na + ] low, [K + ] high; outside cell [Na + ] high, [K + ] low. Outer cell membrane is positive.

On nerve stimulation, Na + enters cell, inner membrane becomes more positive, then K + ions leave cell to reestablish positive potential on outside. These changes occur on a millisecond timescale.

13 ATP Synthesis So far looked at electrons, but what about the protons? NADH + H + NAD + + 2H + + 2e - As redox processes occur, free energy used to force protons into the intermembrane space.

5 Concentration Cells Nerve cells operate as concentration cells. Inside cell [Na + ] low, [K + ] high; outside cell [Na + ] high, [K + ] low. Outer cell membrane is positive. One third of our ATP is used to maintain this difference. On nerve stimulation, Na + enters cell, inner membrane becomes more positive, then K + ions leave cell to reestablish positive potential on outside. These changes occur on a millisecond timescale. Large changes in charge in one region of the membrane stimulate the neighbouring region and the impulse moves down the length of the cell. 13 ATP Synthesis So far looked at electrons, but what about the protons? NADH + H + NAD + + 2H + + 2e - As redox processes occur, free energy used to force protons into the intermembrane space. This creates a concentration cell across the membrane. When [H + ] intermembrane /[H + ] matrix ~ 2.5 a trigger allows protons to flow back across membrane and this spontaneous process drives the non-spontaneous formation of ATP. 14 Chemical Kinetics The rate of a reaction is the speed with which the concentrations of the molecules present change. The rate is given by the gradient of a concentration vs time graph. 15 5

6 Decomposition of H 2 O 2 2H 2 O 2 (aq) 2H 2 O + O 2 (g) Start M H 2 O 2 After 10 s [H 2 O 2 ] = M [H 2 O 2 ]= M Time interval t = s Rate of reaction = Rate of change of [H 2 O 2 ] = [ H 2 O 2 ]/ t = /10.0 = -1.7 x 10-3 M s -1 To avoid negative rates. Rate is defined as - [reactant]/ t 16 Decomposition of H 2 O 2 The rate of removal of H 2 O 2 is not constant. The lower the concentration of H 2 O 2 the slower the rate. The rate is dependent on concentration. Concentraion high rate low rate Time 2H 2 O 2 (aq) 2H 2 O (l) + O 2 (g) Time s t s [H2O2] [H2O2] M M Reaction rate x 10-4 M s The rate of reaction The rate of reaction depends on (among other factors)... Concentration of some or all of the molecules present Temperature The presence of a catalyst The rate equation can only be determined by experiment, not from the stoichiometric equation 18 6

7 The Rate Equation This relates the rate of reaction to the concentration terms which affect it. For the general reaction: A + B M + N -d[a]/dt = k [A] x [B] y Rate constant, k, is a characteristic of the reaction and depends on temperature. Order of the reaction is given by x + y. x is the order with respect to A, y is the order with respect to B etc. 19 Example H 2 (g) + 2 ICl (g) 2 HCl (g) + I 2 (s) Rate = -d[h 2 ] / dt= k[h 2 ][ICl] (by experiment) This reaction is first order with respect to H 2, first order with respect to ICl and second order overall. Question: For the reaction: 2NO 2 (g) + F 2 (g) 2NO 2 F (g) Rate = k[no 2 ][F 2 ] What is the order of reaction with respect to NO 2, F 2 and the overall order of reaction? 20 Determine the Rate Equation O 2 (g) + 2NO(g) 2NO 2 (g) [O 2 ] / M [NO] / M Initial Rate / M s x x x x x x x x x x x x 10-3 Rate = k [O 2 ] x [NO] y keep one conc constant, alter other conc and see what happens to rate. 21 7

8 Determine the Rate Equation O 2 (g) + 2NO(g) 2NO 2 (g) [O 2 ] / M [NO] / M Initial Rate / M s x x x x x x x x x x x x 10-3 Rate = k [O 2 ] x [NO] y double double therefore x = 1 (2 = (2) 1 ) 22 Determine the Rate Equation O 2 (g) + 2NO(g) 2NO 2 (g) [O 2 ] / M [NO] / M Initial Rate / M s x x x x x x x x x x x x 10-3 Rate = k [O 2 ] 1 [NO] y four times double therefore y = 2 (4 = (2) 2 ) 23 Determine the Rate Equation O 2 (g) + 2NO(g) 2NO 2 (g) [O 2 ] / M [NO] / M Initial Rate / M s x x x x x x x x x x x x 10-3 Rate = k [O 2 ] 1 [NO] 2 substitute values to get k 3.21 x 10-3 = k (1.10 x 10-2 )(1.30 x 10-2 ) 2 Rate constant, k = 1.73 x 10 3 s -1 M

Question: Determine the rate equation and value of the rate constant for the reaction: NO 2 (g) + CO (g) NO (g) + CO 2 (g) [NO 2 ]/M [CO] / M Initial Rate / M s -1 1 0.10 0.10 0.0050 2 0.40 0.10 0.080 3 0.

9 Question: Determine the rate equation and value of the rate constant for the reaction: NO 2 (g) + CO (g) NO (g) + CO 2 (g) [NO 2 ]/M [CO] / M Initial Rate / M s Half Life, t 1/2 The half life of a reaction is the time required for the concentration to fall to half its initial value. For first order reactions this is a constant. First Order Reactions t 1/2 = Ln 2 / k and Ln[A]=Ln[A] o -k t 26 Example An antibiotic breaks down in the body with a first order rate constant of k = 1.9 x 10-2 min -1. How long does it take for the concentration to drop to 10% of the initial value? rearrange Ln[A]=Ln[A] o -kt to give Ln( [A]/[A] o )= -k t We do not know the actual concentrations but can use the ratio of concentrations: Ln (10/100) = -(1.9 x 10-2 ) t Ln 0.1 = = -(1.9 x 10-2 ) t and t = / 1.9 x 10-2 =120 min 27 9

10 Question The radioactive isotope 15 O is used in medical imaging and has a half life of seconds. What percentage is left up after 10 minutes? Use k =Ln2/t ½ and Ln[A]=Ln[A] o -kt Note: Ln{ [A]/[A] o }= -k t and [A]/[A] o is amount left. 28 Reaction Mechanisms A balanced chemical equation describes the overall chemical reaction e.g. 2 NO 2 (g) + F 2 (g) 2 NO 2 F (g) A reaction mechanism is a series of elementary reaction steps that add up to give a detailed description of a chemical reaction. e.g. NO 2 + F 2 NO 2 F + F Step 1 NO 2 + F NO 2 F Step 2 2 NO 2 + F 2 2 NO 2 F Overall 29 Reaction Mechanisms A rate determining step in a reaction mechanism is an elementary process that is the slowest step in the mechanism e.g. Step 1 NO 2 + F 2 NO 2 F + F (slow) In simple reactions, the exponents in the rate equation are the same as the coefficients of the molecules of the rate determining elementary process. So in this case, Rate = k [NO 2 ][F 2 ] Note the overall reaction is: 2NO 2 (g) + F 2 (g) 2NO 2 F(g) 30 10

Arrhenius Equation k = Ae -E a/rt The activation energy, E a, of a reaction is the minimum amount of energy that the reacting molecules must possess if

The Arrhenius equation describes the temperature dependence of the rate constant that is exponentially related to the activation energy (A is the

11 Arrhenius Equation k = Ae -E a/rt The activation energy, E a, of a reaction is the minimum amount of energy that the reacting molecules must possess if the reaction is to be successful. The Arrhenius equation describes the temperature dependence of the rate constant that is exponentially related to the activation energy (A is the pre-exponential factor, or the A factor ). Energy activation energy exothermic reaction Reaction coordinate 31 Reaction Profile 32 Arrhenius Equation k = Ae -E a/rt At higher temperatures, more molecules have the minimum required energy (E a ) and hence the faster the reaction

Catalysis A catalyst provides an alternative reaction

In this case more molecules have the minimum energy

12 Catalysis A catalyst provides an alternative reaction pathway of lower activation energy. In this case more molecules have the minimum energy required for successful reaction and the reaction proceeds at a faster rate. 34 Enzyme Catalysis Efficient: rate enhancements of 10 8 to possible. Specific. But low tolerance to temperature and ph changes

17.1 Redox Chemistry Revisited

Chapter Outline 17.1 Redox Chemistry Revisited 17.2 Electrochemical Cells 17.3 Standard Potentials 17.4 Chemical Energy and Electrical Work 17.5 A Reference Point: The Standard Hydrogen Electrode 17.6